1. Oxidation Reduction

2. Definitions ► Oxidation- an element has lost electrons to another element. ► Reduction- an element gains electrons from another element. ► Oxidizing agent- a substance which causes another substance to be oxidized. itself is reduced. ► Oxidizer- same as oxidizing agent. ► Reducing Agent- a substance which causes another substance to be reduced. Itself becomes oxidized. ► Reducer or reductant- same as above. ► Anion- (-) ion. ► Cation- (+) ion ►.►.►.►. Cathode: positive charge, where oxidation occursCathode: positive charge, where oxidation occurs

3. ► Electrolyte- The salt solution in a battery or electrochemical cell. ► Electrodes- are in batteries or an electrochemical cell. Location for either reduction or oxidation.  Anode: charged electrode where oxidation occurs.  Cathode: charged electrode where reduction occurs.

4. Rules for determining the oxidation number(state) of a substance. ► Page 159 in Zumdahl ► 1. The oxidation state of an atom in its elemental state is 0. For ex:, the oxidation state of each atom in the substances Na(s), O 2 (g), O 3 (g), and Hg(l) is 0. ► 2. The oxidation state of a monatomic ion is the same as its charge. For example, the oxidation state of the Na ion is +1 and of the Cl ion is -1

5. Rules continued… ► 3. Oxygen is assigned an oxidation state of -2 in its covalent compounds such as CO, CO 2, SO 2, and SO 3. An exception to this rule occurs in peroxides( compounds containing the O 2 group), where each oxygen is assigned an oxidation state of -1. The best known example of a peroxide is hydrogen peroxide ( H 2 O 2 )

6. Rules continued… ► 4. In its covalent compounds with nonmetals, hydrogen is assigned an oxidation state of +1. For ex: in the compounds HCl, NH 3, H 2 O, and CH 4, hydrogen is assigned an oxidation state of +1. ► 5. In its compounds, fluorine is always assigned an oxidation state of -1.

7. Rules continued…. ► 6. The sum of the oxidation states must be zero for an electrically neutral compound. For an ion, the sum of the oxidation states must equal the charge of the ion. For ex: the sum of the oxidation states for the hydrogen and oxygen atoms in water is 0; the sum of the oxidation states for the carbon and oxygen atoms in CO 3 is -2; and the sum of oxidation states for the nitrogen and hydrogen atoms in NH 4 + is +1 ► 7. Group 1 and 2 metals: oxidation state = ionic charge.

8. Determining an oxidation number Examples: K 2 Cr 2 O 7 CH 4 H2H2H2H2 Br – FeO Fe 2 O 3 C 2 H 5 OH CuCl 2

9. Determining what was oxidized and what was reduced ► Examples: ► Mg + 2HCl  MgCl 2 + H 2 ► N 2 + 3H 2  2NH 3 ► CH 4 + 2O 2  CO 2 + 2H 2 O ► 2Na + 2H 2 O  NaOH + H 2 ► Fe 2 O 3 + 3CO  2Fe + 3CO 2 ► Na 2 SO 4 + BaCl 2  BaSO 4 + 2NaCl

10. Relative reactivities ► Ultimate goal:  To be able to predict whether a reaction will occur or not (be spontaneous or not)  To predict the products produced in a spontaneous reaction  To be able to determine relative reactivities for a series of elements.

11. Determining if a reaction is spontaneous and the products  A reactivity chart is used  The more easily reduced substance must end up in its reduced state, the other substance must end up in its oxidized state.  If the substances are already in their most reduced or oxidized state than no reaction occurs.

12. continued ► Oxidized state: ► For metals:  Cation state: ex. Mg +2, Na +,  transition metals must have their highest oxidation number. Ex Fe +3, not Fe +2 ► For nonmetals:  Diatomic molecules:ex F 2, O 2, Cl 2,, etc. ► Reduced state: ► For metals:  Solid state: ex. Mg, Na, Fe etc. ► For nonmetals:  Anion state: ex. F-, Cl-, O -2, N -3 etc.

13. continued ► Hydrogen, even though it is a nonmetal, has these states:  Oxidized state: H +  Reduced state: H 2

14. Determining whether a reaction is spontaneous ► Examples:  Cu(s) + Zn +2   Cu +2 + Zn(s)   I 2 + 2Cl -   2I - + Cl 2   Zn + H +   Cu + H + 

15. Determining relative reactivities ► Example  Rank the following elements from strongest reducing agent to weakest reducing agent.  A + B +2  no reaction  A + C +2  A +2 + C  C + D +2  no reaction  A + D +2  A +2 + D

16. ► Example: ► rank the elements from strongest oxidizer to weakest oxidizer.  Zn +2 + Cu  no reaction  Zn +2 + Mg  Zn + Mg +2  Zn +2 + Pb  no reaction  Pb + Cu +2  Pb +2 + Cu

17. The two electrochemical cells ► Battery or voltaic cell  Uses a spontaneous redox reaction  Produces electricity ► Electrolytic cell  Called electrolysis  USES A NONSPONTANEOUS REDOX reaction  Uses electricity  Practical use is to obtain elements in a pure reduced form that cannot be obtained in another way.

18. Battery: Voltaic cell or galvanic cell ► Produces electricity ► Must use a spontaneous redox reaction ► Oxidation occurs at the anode (-) ► Reduction occurs at the cathode (+)

19. Voltaic cell ► Anode  Is negatively charged  Oxidation occurs  The more easily oxidized element is located here.  e- are removed from the neutral metal.  Ex. Zn (s)  Zn +2 + 2e- ► Cathode  Is positively charged  Reduction occurs  The more easily reduced metal is located here.  e- are added to the cation of the metal to produce a pure elemental substance.  2e- + Cu +2  Cu (s)

20. A battery with zinc and copper ► Zinc will spontaneously lose e- to the copper. ► Anode:  Solid piece of Zn metal  Zn has e- removed from it.  The e- move through the wire to the cathode.  Zinc ions are formed and are in the solution.

21. ► Cathode:  Copper ions are gaining e- and being reduced.  Solid Cu forms around the outside of the Cu cathode.  Other info: the electrolyte is needed to complete the circuit, and must be free moving.  There must be a barrier between the electrodes or the copper will get e- directly from the zinc and the e- will not move through the wires.

22. How a voltaic cell works

24. Predicting the battery arrangement between lead and silver ► For a reaction between lead and silver:  Determine what will be oxidized and what will be reduced  Draw a diagram of a simple battery. Label the anode and the cathode including its corresponding charge.  Write out the half reaction that occurs at each electrode.  Write the over all redox reaction that occurs.

26. Diagram of a battery

27. Electrolytic cell ► Cathode ► Negative charge ► Attracts the cation ► Electrons are given to the cation ► Reduction occurs ► Ex: Na + + 1e -  Na (s) ► A build up of a fuzzy metal, or a gas appears ► Anode ► Positive charge ► Attracts anions ► Electrons are removed from the anions. Neutral elements are formed ► Oxidation occurs ► 2Cl-  Cl 2 + 2e- ► A gas appears

28. Electrolytic cell reactions ► When NaCl is used:  2(Na + + 1e -  Na) reduction  2Cl -  Cl 2 + 2e - oxidation  note: the electrons must cancel out. There must be the same # of e - in reduction as there is an oxidation.

29. continued ► When water undergoes electrolysis  2H 2 O  2H 2 + O 2  2(2H + + 2e -  H 2 ) reduction  2O -2  O 2 + 4e - oxidation

30. Electolytic cell: electrolyisis of NaCl